[kathhhriine]

Hi.

I'm currently learning about enthalpy and thermochemistry, and I've done quite a few exercises about this. However, I keep using the wrong enthalpy formulas. I've so far learnt that there is two types of enthalpy; enthalpy of combustion and enthalpy of formation, and to calculate these I use these formulas: ΔH_reaction=H_combustion(reactants)-H_combustion(products), and ΔH_reaction=H_formation(products)-H_formation(reactants).
Take for an instant this question:
    N2(g) + O2(g) -> NO2(g)   ΔH^Ө = –57 kJ mol–1
       N2(g) + 2O2(g) -> N2O4(g)   ΔH^Ө = +9 kJ mol–1
 calculate ΔH^Ө for the following reaction (in kJ): 2NO2(g) -> N2O4(g)

Here I thought Id use the formation enthalpy formula, getting 9- (2x57)=-105kJmol^-1, however this was wrong, and the answer says it is +123kJmol^-1, implying that the formula for combustion is used.

Could someone please explain why this is, and when to apply the different formulas?

[mjc123]

You have made a sign mistake; it is +9 - (2*-57) = 9+114 = +123
And your first equation is unbalanced.

[kathhhriine]

You have made a sign mistake; it is +9 - (2*-57) = 9+114 = +123
And your first equation is unbalanced.

Thank you! This clear up the answers for a lot of the exercises I've done
Do you have an explanation for when to use ΔH=H(prod)-H(react) and ΔH=H(react)-H(prod) ?

[mjc123]

A lot? Don't you know how to manipulate signs?

Consider the schemes
A. Reactants  Elements  Products
B. Reactants + O₂  Oxidation products  Products + O₂
In each case, can you work out an expression for ΔH(reactants  products) in terms of enthalpies of formation or combustion?

Tip: stop asking what formula to use. Start thinking logically about the physical situation.